Bronsted-Lowry Theory

Learning Objectives (OCR Spec 5.1.3)

By the end of this lesson you should be able to:

• State the Bronsted-Lowry definition of acids and bases
• Identify conjugate acid-base pairs in reactions
• Distinguish between monobasic, dibasic and tribasic acids
• Write equations showing proton transfer between acid and base
• Recognise the role of water as either an acid or a base (amphoteric behaviour)

The Bronsted-Lowry Definition

Early definitions of acids (sour taste, turning litmus red) were inadequate because they described properties, not structure. In 1923, Johannes Bronsted and Thomas Lowry independently proposed a definition based on proton transfer:

• An acid is a proton (H+) donor.
• A base is a proton (H+) acceptor.

Every acid-base reaction is therefore a proton transfer from the acid to the base:

HA + B -> A- + BH+

The acid HA loses a proton to become A-. The base B gains a proton to become BH+. Notice that the reaction is reversible in principle - the products can themselves act as a (weaker) acid and base.

Worked Example: HCl and Water

When hydrogen chloride dissolves in water the following equilibrium is set up (for HCl it lies far to the right):

HCl(aq) + H2O(l) -> H3O+(aq) + Cl-(aq)

• HCl donates a proton -> HCl is the acid
• H2O accepts a proton -> H2O is the base
• H3O+ is the oxonium (hydronium) ion

H3O+ is the species that gives acidic solutions their properties. At A-Level we often simplify H3O+(aq) to H+(aq) for clarity, but you should recognise that the bare proton does not exist free in water; it is always solvated.

Conjugate Acid-Base Pairs

When an acid donates a proton, the species left behind can accept a proton back. This species is called the conjugate base of the acid. Likewise, when a base accepts a proton, the resulting species can donate it back - this is the conjugate acid of the base.

In HCl + H2O -> H3O+ + Cl-:

Role	Species	Conjugate
Acid 1	HCl	Cl- (conjugate base of HCl)
Base 1	H2O	H3O+ (conjugate acid of H2O)

Two conjugate acid-base pairs appear in every Bronsted-Lowry reaction. They differ by a single proton:

graph LR
  A[HCl<br/>acid 1] -->|loses H+| B[Cl-<br/>base 1]
  C[H2O<br/>base 2] -->|gains H+| D[H3O+<br/>acid 2]

The pairs are HCl / Cl- and H2O / H3O+. Each pair differs by exactly one H+.

Another Example: Ammonia in Water

NH3(aq) + H2O(l) <=> NH4+(aq) + OH-(aq)

• NH3 accepts a proton -> NH3 is the base, NH4+ is its conjugate acid
• H2O donates a proton -> H2O is the acid, OH- is its conjugate base

Ammonia is a weak base, so the equilibrium lies to the left - the solution contains mostly NH3 with a small amount of NH4+ and OH-.

Water is Amphoteric

Notice that in the HCl example water behaves as a base (accepts a proton) but in the NH3 example water behaves as an acid (donates a proton). A species that can act as either a Bronsted-Lowry acid or a Bronsted-Lowry base is said to be amphoteric.

Water is the archetypal amphoteric solvent:

• H2O + HCl -> H3O+ + Cl- (water is base)
• H2O + NH3 <=> OH- + NH4+ (water is acid)

This dual behaviour underlies the self-ionisation of water (lesson 2, Kw).

Monobasic, Dibasic and Tribasic Acids

Some acids can release more than one proton per molecule. The basicity of an acid is the number of replaceable hydrogen atoms (protons it can donate).

Acid	Formula	Basicity	Ionisation steps
Hydrochloric	HCl	Monobasic (monoprotic)	HCl -> H+ + Cl-
Nitric	HNO3	Monobasic	HNO3 -> H+ + NO3-
Ethanoic	CH3COOH	Monobasic	CH3COOH <=> H+ + CH3COO-
Sulfuric	H2SO4	Dibasic (diprotic)	H2SO4 -> H+ + HSO4-; HSO4- <=> H+ + SO4^2-
Carbonic	H2CO3	Dibasic	H2CO3 <=> H+ + HCO3-; HCO3- <=> H+ + CO3^2-
Phosphoric	H3PO4	Tribasic (triprotic)	three successive ionisations

Exam tip: not every H atom in an acid molecule is acidic. In ethanoic acid CH3COOH only the COOH hydrogen is acidic; the three CH3 hydrogens are not donated. The basicity is therefore 1, not 4. Acidic hydrogens are those attached to strongly electronegative atoms (O, halogens) where the O-H (or similar) bond is polar enough to release H+ in water.

Reactions of Acids (Bronsted-Lowry View)

All the familiar acid reactions can be seen as proton transfer from the acid to a base:

Acid + Metal

Mg(s) + 2HCl(aq) -> MgCl2(aq) + H2(g)

Here the metal is not a Bronsted-Lowry base; this is a redox reaction rather than a pure acid-base reaction. But the acid still behaves as a proton donor at the metal surface.

Acid + Base (Alkali)

HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l) Net ionic: H+(aq) + OH-(aq) -> H2O(l)

OH- accepts a proton to form water. This is the archetypal neutralisation.

Acid + Carbonate / Hydrogencarbonate

2HCl(aq) + Na2CO3(aq) -> 2NaCl(aq) + H2O(l) + CO2(g) HCl(aq) + NaHCO3(aq) -> NaCl(aq) + H2O(l) + CO2(g)

The carbonate or hydrogencarbonate ion is the base. It accepts protons, the resulting H2CO3 decomposes rapidly to H2O and CO2, which is why effervescence is observed.

Acid + Ammonia

HCl(aq) + NH3(aq) -> NH4Cl(aq) Net ionic: H+(aq) + NH3(aq) -> NH4+(aq)

NH3 is the base.

In every case one species donates a proton and another accepts it. Identifying the pairs is the first step to writing any acid-base equation correctly.

Worked Example: Identifying Pairs

Identify the conjugate acid-base pairs in:

HSO4-(aq) + H2O(l) <=> SO4^2-(aq) + H3O+(aq)

Solution:

• HSO4- loses a proton to become SO4^2-, so HSO4- is acid 1 and SO4^2- is its conjugate base.
• H2O gains a proton to become H3O+, so H2O is base 2 and H3O+ is its conjugate acid.

Pairs: HSO4- / SO4^2- and H2O / H3O+.

Note that HSO4- itself is also the conjugate base of H2SO4. The same species can be an acid in one reaction and a base in another - it depends on the partner.

Exam Tips

• Always start by asking: which species loses a proton, and which gains one? That identifies the acid and base.
• Write H3O+ for the aqueous proton if the question asks for precise Bronsted-Lowry species; otherwise H+(aq) is acceptable.
• For conjugate pairs, they must differ by a single H+. If you have added or removed anything else you have misidentified the pair.
• Remember that basicity counts only the hydrogens that can actually be donated, not the total number of H atoms in the formula.
• Use <=> for weak acid ionisations and -> for strong acid ionisations (lessons 3-5).
• Water is amphoteric - do not be surprised to see it on either side as acid or base.

Common Exam Mistakes

• Calling a lone pair donor (e.g. NH3) a "proton donor" - it is a proton acceptor, i.e. a base.
• Writing H+ alone as an acid - H+ is what is being transferred; the acid is the species that gives it up.
• Mixing up "conjugate" with "complementary" or "opposite"; the conjugate base is simply what the acid becomes after losing H+.
• Miscounting basicity because of non-acidic hydrogens (as in CH3COOH).
• Forgetting that some reactions are written -> (complete) and others <=> (equilibrium) - this matters hugely for strong vs weak acids later on.

Summary

Bronsted-Lowry theory defines acid-base chemistry in terms of proton transfer: acids donate H+, bases accept H+. Each reaction contains two conjugate acid-base pairs, each differing by one H+. Water is amphoteric, acting as either acid or base depending on the partner. Basicity (mono-, di-, tri-) counts only the protons an acid can actually donate. These ideas are the foundation for quantitative pH, Ka and buffer calculations in the rest of this module.

Further Worked Examples

Worked Example: Ammonium ion as acid

Identify the acid, base and conjugates in:

NH4+(aq) + H2O(l) <=> H3O+(aq) + NH3(aq)

Solution: NH4+ donates a proton -> acid; H2O accepts a proton -> base. Pairs: NH4+ / NH3 and H2O / H3O+. Notice this is the opposite of the NH3 + H2O example earlier: here the ammonium ion acts as the acid because the proton is transferred from NH4+ to water.

Worked Example: Carbonate with acid

Write the Bronsted-Lowry equation for the first step of the reaction between carbonate ion and ethanoic acid.

CH3COOH(aq) + CO3^2-(aq) <=> CH3COO-(aq) + HCO3-(aq)

Here CH3COOH is acid 1 (donates H+) and CO3^2- is base 2 (accepts H+). The conjugate pairs are CH3COOH / CH3COO- and HCO3- / CO3^2-. A second proton transfer can then produce H2CO3, which decomposes to CO2 and H2O - hence the familiar effervescence.

A Brief Historical Note

Before Bronsted and Lowry, Svante Arrhenius defined acids as species that dissociate in water to give H+ ions, and bases as species that dissociate to give OH-. This works for HCl, HNO3 and NaOH but fails to explain why NH3 (with no OH- in its formula) is a base, or why reactions can occur in non-aqueous solvents. The Bronsted-Lowry extension replaced "gives H+/OH- in water" with "donates/accepts protons" - a broader definition that works in any solvent and covers ammonia as a base. An even more general definition was later proposed by G. N. Lewis (acids as electron-pair acceptors), but at A-Level the Bronsted-Lowry view is usually sufficient.

In the next lesson we will meet the ionic product of water Kw, which turns the qualitative idea of water self-ionising into a quantitative equation we can use to calculate pH.